Chemistry Notes for Class 11- Structure of Atom-Rutherford’s Atomic Model
Chemistry notes for class 11: Rutherford’s Atomic Model
Ernest Rutherford and his co-workers were working in the area of radioactivity. They were studying the effect of alpha (α) particles on the matter. The alpha particles are helium nuclei, which can be obtained by the removal of two electrons from the helium atom. In 1910, Hans Geiger (Rutherford’s technician) and Ernest Marsden (Rutherford’s student) performed the famous α-ray scattering experiment. This led to the failure of Thomson’s model of the atom. Let us learn about this experiment.
α-Ray scattering experiment
In this experiment, a stream of α particle from a radioactive source was directed on a thin (about 0.00004 cm thick) piece of gold foil. According to Thomson’s model, it was expected that the alpha particles would just pass straight through the gold foil and could be detected by a photographic plate placed behind the foil. But the actual results of the experiment were quite surprising. It was observed that:
(i) Most of the α-particles passed straight through the gold foil.
(ii) Some of the α-particles were deflected by small angles.
(iii) A few particles were deflected by large angles.
(iv) About 1 in every 12000 particles experienced a rebound.
The results of α-ray scattering experiment were explained by Rutherford in 1911 and another model of the atom was proposed. According to Rutherford’s model, an atom contains a dense and positively charged region located at its centre; it was called as the nucleus, all the positive charge of an atom and most of its mass was contained in the nucleus. The rest of an atom must be empty space which contains the much smaller and negatively charged electrons.
Schematic Figure shows Rutherford’s Experiment
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On the basis of the proposed model, the experimental observations in the scattering experiment could be explained. The α particles passing through the atom in the region of the electrons would pass straight without any deflection. Only those particles that come in the close vicinity of the positively charged nucleus get deviated from their path. Very few α-particles, those that collide with the nucleus, would face a rebound.On the basis of his model, Rutherford was able to predict the size of the nucleus.He estimated that the radius of the nucleus was at least 1/10000 times smaller than that of the radius of the atom.
Based on his observations, Rutherford proposed the following structural features of an atom:
- Most of the atom’s mass and its entire positive charge are confined to a small core, called nucleus. The positively charged particle is called proton.
- Most of the volume of an atom is empty space.
- The number of negatively charged electrons dispersed outside the nucleus is same as a number of positively charge in the nucleus. It explains the overall electrical neutrality of an atom.
Rutherford’s Model of an Atom
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Limitations of Rutherford’s model
According to Rutherford’s model, the negatively charged electrons revolve in circular orbits around the positively charged nucleus. However, according to Maxwell’s electromagnetic theory, if a charged particle accelerates around another charged particle then it would continuously lose energy in the form of radiation. This suggests that loss of energy would slow down the speed of the electron. Therefore, the electron is expected to move in a spiral fashion around the nucleus and eventually fall into the nucleus.
In Rutherford’s model of the atom, an accelerating electron should gradually lose energy and eventually spiral into the positively charged nucleus.
In other words, the atom must not be stable. However, we know that the atom is stable and such a collapse does not occur. Thus, Rutherford’s model is unable to explain the stability of the atom. We know that an atom may contain a number of electrons. The Rutherford’s model also does not say anything about the way the electrons are distributed around the nucleus. Another drawback of Rutherford’s model was its inability to explain the relationship between the atomic mass and atomic number (the number of protons).
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